Common Structures of Binary Compounds

As shown in part (a) in Figure 12.8 "Holes in Cubic Lattices", a simple cubic lattice of anions contains only one kind of hole, located in the center of the unit cell. Because this hole is equidistant from all eight atoms at the corners of the unit cell, it is called a cubic holeThe hole located at the center of the simple cubic lattice. The hole is equidistant from all eight atoms or ions at the corners of the unit cell. An atom or ion in a cubic hole has a coordination number of 8.. An atom or ion in a cubic hole therefore has a coordination number of 8. Many ionic compounds with relatively large cations and a 1:1 cation:anion ratio have this structure, which is called the cesium chloride structureThe unit cell for many ionic compounds with relatively large cations and a 1:1 cation:anion ratio. (Figure 12.9 "The Cesium Chloride Structure") because CsCl is a common example.Solid-state chemists tend to describe the structures of new compounds in terms of the structure of a well-known reference compound. Hence you will often read statements such as “Compound X possesses the cesium chloride (or sodium chloride, etc.) structure” to describe the structure of compound X. Notice in Figure 12.9 "The Cesium Chloride Structure" that the z = 0 and the z = 1.0 planes are always the same. This is because the z = 1.0 plane of one unit cell becomes the z = 0 plane of the succeeding one. The unit cell in CsCl contains a single Cs⁺ ion as well as $8\times \frac{1}{8}{\text{Cl}}^{-}=1{\text{Cl}}^{-}$ ion, for an overall stoichiometry of CsCl. The cesium chloride structure is most common for ionic substances with relatively large cations, in which the ratio of the radius of the cation to the radius of the anion is in the range shown in Table 12.2 "Relationship between the Cation:Anion Radius Ratio and the Site Occupied by the Cations".

Figure 12.8 Holes in Cubic Lattices

The three illustrations show (a) the cubic hole that is in the center of a simple cubic lattice of anions, (b) the locations of the octahedral holes in a face-centered cubic lattice of anions, and (c) the locations of the tetrahedral holes in a face-centered cubic lattice of anions.

Figure 12.9 The Cesium Chloride Structure

The Cs⁺ ion occupies the cubic hole in the center of a cube of Cl⁻ ions. The drawings at the right are horizontal cross-sections through the unit cell at the bottom (z = 0) and halfway between the bottom and top (z = 0.5). A top cross-section (z = 1) is identical to z = 0. Such cross-sections often help us visualize the arrangement of atoms or ions in the unit cell more easily.

In contrast, a face-centered cubic (fcc) array of atoms or anions contains two types of holes: octahedral holesOne of two kinds of holes in a face-centered cubic array of atoms or ions (the other is a tetrahedral hole). One octahedral hole is located in the center of the face-centered cubic unit cell, and there is a shared one in the middle of each edge. An atom or ion in an octahedral hole has a coordination number of 6., one in the center of the unit cell plus a shared one in the middle of each edge (part (b) in Figure 12.8 "Holes in Cubic Lattices"), and tetrahedral holesOne of two kinds of holes in a face-centered cubic array of atoms or ions (the other is an octahedral hole). Tetrahedral holes are located between an atom at a corner and the three atoms at the centers of the adjacent faces of the face-centered cubic unit cell. An atom or ion in a tetrahedral hole has a coordination number of 4., located between an atom at a corner and the three atoms at the centers of the adjacent faces (part (c) in Figure 12.8 "Holes in Cubic Lattices"). As shown in Table 12.2 "Relationship between the Cation:Anion Radius Ratio and the Site Occupied by the Cations", the ratio of the radius of the cation to the radius of the anion is the most important determinant of whether cations occupy the cubic holes in a cubic anion lattice or the octahedral or tetrahedral holes in an fcc lattice of anions. Very large cations occupy cubic holes in a cubic anion lattice, cations of intermediate size tend to occupy the octahedral holes in an fcc anion lattice, and relatively small cations tend to occupy the tetrahedral holes in an fcc anion lattice. In general, larger cations have higher coordination numbers than small cations.

The most common structure based on a fcc lattice is the sodium chloride structureThe solid structure that results when the octahedral holes of an fcc lattice of anions are filled with cations. (Figure 12.10 "The Sodium Chloride Structure"), which contains an fcc array of Cl⁻ ions with Na⁺ ions in all the octahedral holes. We can understand the sodium chloride structure by recognizing that filling all the octahedral holes in an fcc lattice of Cl⁻ ions with Na⁺ ions gives a total of 4 Cl⁻ ions (one on each face gives $6\times \frac{1}{2}=3$ plus one on each corner gives $8\times \frac{1}{8}=1$, for a total of 4) and 4 Na⁺ ions (one on each edge gives $12\times \frac{1}{4}=3$ plus one in the middle, for a total of 4). The result is an electrically neutral unit cell and a stoichiometry of NaCl. As shown in Figure 12.10 "The Sodium Chloride Structure", the Na⁺ ions in the sodium chloride structure also form an fcc lattice. The sodium chloride structure is favored for substances with two atoms or ions in a 1:1 ratio and in which the ratio of the radius of the cation to the radius of the anion is between 0.414 and 0.732. It is observed in many compounds, including MgO and TiC.

Figure 12.10 The Sodium Chloride Structure

In NaCl, the Na⁺ ions occupy the octahedral holes in an fcc lattice of Cl⁻ ions, resulting in an fcc array of Na⁺ ions as well.

The structure shown in Figure 12.11 "The Zinc Blende Structure" is called the zinc blende structureThe solid structure that results when half of the tetrahedral holes in an fcc lattice of anions are filled with cations with a 1:1 cation:anion ratio and a coordination number of 4., from the common name of the mineral ZnS. It results when the cation in a substance with a 1:1 cation:anion ratio is much smaller than the anion (if the cation:anion radius ratio is less than about 0.414). For example, ZnS contains an fcc lattice of S²⁻ ions, and the cation:anion radius ratio is only about 0.40, so we predict that the cation would occupy either a tetrahedral hole or an octahedral hole. In fact, the relatively small Zn²⁺ cations occupy the tetrahedral holes in the lattice. If all 8 tetrahedral holes in the unit cell were occupied by Zn²⁺ ions, however, the unit cell would contain 4 S²⁻ and 8 Zn²⁺ ions, giving a formula of Zn₂S and a net charge of +4 per unit cell. Consequently, the Zn²⁺ ions occupy every other tetrahedral hole, as shown in Figure 12.11 "The Zinc Blende Structure", giving a total of 4 Zn²⁺ and 4 S²⁻ ions per unit cell and a formula of ZnS. The zinc blende structure results in a coordination number of 4 for each Zn²⁺ ion and a tetrahedral arrangement of the four S²⁻ ions around each Zn²⁺ ion.

Figure 12.11 The Zinc Blende Structure

Zn²⁺ ions occupy every other tetrahedral hole in the fcc array of S²⁻ ions. Each Zn²⁺ ion is surrounded by four S²⁻ ions in a tetrahedral arrangement.

We examine only one other structure of the many that are known, the perovskite structureA structure that consists of a bcc array of two metal ions, with one set (M) located at the corners of the cube, and the other set (M′) in the centers of the cube.. Perovskite is the generic name for oxides with two different kinds of metal and have the general formula MM′O₃, such as CaTiO₃. The structure is a body-centered cubic (bcc) array of two metal ions, with one M (Ca in this case) located at the corners of the cube, and the other M′ (in this case Ti) in the centers of the cube. The oxides are in the centers of the square faces (part (a) in Figure 12.12 "The Perovskite Structure of CaTiO"). The stoichiometry predicted from the unit cell shown in part (a) in Figure 12.12 "The Perovskite Structure of CaTiO" agrees with the general formula; each unit cell contains $8\times \frac{1}{8}=1$ Ca, 1 Ti, and $6\times \frac{1}{2}=3$ O atoms. The Ti and Ca atoms have coordination numbers of 6 and 12, respectively. We will return to the perovskite structure when we discuss high-temperature superconductors in Section 12.7 "Superconductors".

Figure 12.12 The Perovskite Structure of CaTiO₃

Two equivalent views are shown: (a) a view with the Ti atom at the center and (b) an alternative view with the Ca atom at the center.

X-Ray Diffraction

As you learned in Chapter 6 "The Structure of Atoms", the wavelengths of x-rays are approximately the same magnitude as the distances between atoms in molecules or ions. Consequently, x-rays are a useful tool for obtaining information about the structures of crystalline substances. In a technique called x-ray diffractionAn technique used to obtain information about the structures of crystalline substances by using x-rays., a beam of x-rays is aimed at a sample of a crystalline material, and the x-rays are diffracted by layers of atoms in the crystalline lattice (part (a) in Figure 12.13 "X-Ray Diffraction"). When the beam strikes photographic film, it produces an x-ray diffraction pattern, which consists of dark spots on a light background (part (b) in Figure 12.13 "X-Ray Diffraction"). In 1912, the German physicist Max von Laue (1879–1960; Nobel Prize in Physics, 1914) predicted that x-rays should be diffracted by crystals, and his prediction was rapidly confirmed. Within a year, two British physicists, William Henry Bragg (1862–1942) and his son, William Lawrence Bragg (1890–1972), had worked out the mathematics that allows x-ray diffraction to be used to measure interatomic distances in crystals. The Braggs shared the Nobel Prize in Physics in 1915, when the son was only 25 years old. Virtually everything we know today about the detailed structures of solids and molecules in solids is due to the x-ray diffraction technique.

Figure 12.13a X-Ray Diffraction

These illustration shows a schematic drawing of x-ray diffraction.

Figure 12.13b X-Ray Diffraction

This side-by-side comparison shows the X-ray diffraction patterns of two different samples collected from the Martian surface by NASA's Curiosity rover. These images, made from data obtained by Curiosity's Chemistry and Mineralogy instrument (CheMin), show the patterns obtained from a drift of windblown dust and sand called "Rocknest" and from a powdered rock sample drilled from the "John Klein" bedrock. The presence of abundant clay minerals in the John Klein drill powder and the lack of abundant salt suggest a fresh water environment. The presence of calcium sulfates rather than magnesium or iron sulfates (as found at Meridiani Planum by NASA's Mars Exploration Rover Opportunity) suggests a neutral to mildly alkaline pH environment. The Rocknest sand shadow mineralogy suggests a dry, aeolian (wind-shaped) environment with low water activity. The John Klein mineralogy suggests a lacustrine (lakebed) environment with high water activity. As seen on the left, the Rocknest data reveal abundant plagioclase feldspar, pyroxene and olivine minerals. The data also indicate reveal small amounts of magnetite and anhydrite. In addition, the Rocknest sample contains 25 to 35 percent amorphous, or non-crystalline, material. X-ray diffraction analysis of the John Klein drill powder reveals abundant phyllosilicate (a class of clay minerals called smectites that form by the action of relatively pure and neutral pH water on source minerals), plagioclase feldspar, pyroxene, magnetite and olivine. Alternatively, the clay minerals could have been transported by water from sources higher up the sediment fan to form the John Klein mineral assemblage. The region of the pattern indicating the phyllosilicates is labeled in the annotated version of this image. The data also show minor amounts of anhydrite and bassanite. The John Klein sample also contains about 20 percent amorphous material. Image and caption Credit: Public Domain, NASA/JPL-Caltech/Ames, via Wikimedia Commons.

Recall that two waves that are in phase interfere constructively, thus reinforcing each other and generating a wave with a greater amplitude. In contrast, two waves that are out of phase interfere destructively, effectively canceling each other. When x-rays interact with the components of a crystalline lattice, they are scattered by the electron clouds associated with each atom. As shown in Figure 12.5 "The Three Kinds of Cubic Unit Cell", Figure 12.7 "Close-Packed Structures: hcp and ccp", and Figure 12.8 "Holes in Cubic Lattices", the atoms in crystalline solids are typically arranged in planes. Figure 12.14 "The Reflection of X-Rays from Two Adjacent Planes of Atoms Can Result in Constructive Interference of the X-Rays" illustrates how two adjacent planes of atoms can scatter x-rays in a way that results in constructive interference. If two x-rays that are initially in phase are diffracted by two planes of atoms separated by a distance d, the lower beam travels the extra distance indicated by the lines BC and CD. The angle of incidence, designated as θ, is the angle between the x-ray beam and the planes in the crystal. Because BC = CD = d sin θ, the extra distance that the lower beam in Figure 12.14 "The Reflection of X-Rays from Two Adjacent Planes of Atoms Can Result in Constructive Interference of the X-Rays" must travel compared with the upper beam is 2d sin θ. For these two x-rays to arrive at a detector in phase, the extra distance traveled must be an integral multiple n of the wavelength λ:

Equation 12.1 is the Bragg equationThe equation that describes the relationship between two x-ray beams diffracted from different planes of atoms: $2d\sin \theta =n\lambda .$. The structures of crystalline substances with both small molecules and ions or very large biological molecules, with molecular masses in excess of 100,000 amu, can now be determined accurately and routinely using x-ray diffraction and the Bragg equation. Example 4 illustrates how to use the Bragg equation to calculate the distance between planes of atoms in crystals.

Figure 12.14 The Reflection of X-Rays from Two Adjacent Planes of Atoms Can Result in Constructive Interference of the X-Rays

(a) The x-ray diffracted by the lower layer of atoms must travel a distance that is longer by 2d sin θ than the distance traveled by the x-ray diffracted by the upper layer of atoms. Only if this distance (BC plus CD) equals an integral number of wavelengths of the x-rays (i.e., only if λ = 2d sin θ) will the x-rays arrive at the detector in phase. (b) In a solid, many different sets of planes of atoms can diffract x-rays. Each has a different interplanar distance and therefore diffracts the x-rays at a different angle θ, which produces a characteristic pattern of spots.